Acids and Bases

Acids and Bases Introduction

• This lecture serves as a transition into skills and clinical units.
• The lecture content is not included in the case study assessment next week but will be in the final quiz.
• The quiz opens midnight tonight and closes midnight Sunday, including questions on acids and bases.

pH Scale

• Often referred to as the potential of hydrogen.
• Focuses on understanding the behavior of hydrogen concentration.
• Hydrogen (H) is element number one, with one proton in its nucleus and one electron.
• A hydrogen ion (H+) loses its electron and becomes a proton.
  • Very small element and is very important for the human body.
  • Also seen in hydrocarbons.
• Water (H2O) has two hydrogen atoms and one oxygen atom.
• Covalent bonds between hydrogen and oxygen are strong but can break due to the constant movement of water molecules in liquid form.
• When water molecules bump into each other, a hydrogen ion can break away, forming hydroxide (OH-). The electron remains with the Oxygen atom, and creates a negative charge.
• Dissociation: Water molecule splits into hydroxide ion and hydrogen ion.
• Hydrogen as a proton is quickly taken up by other molecules forming H_3O^+, hydronium ion.
• The reaction is reversible, reaching equilibrium in water.
• Pure water has equal concentrations of hydrogen and hydroxide ions, resulting in a neutral pH of 7.
• Solutes that are acids and bases in water disrupt this balance.

pH Scale Measurement

• Measures a solution's acidity or alkalinity.
• Arbitrary logarithmic scale from 0 to 14.
  • 0 is most acidic.
  • 14 is most basic or alkaline.
• pH measures the amount of hydrogen ions in an aqueous solution in moles per liter.
• Acidic solutions have more hydrogen ions (lower than 7 on the pH scale).
• Alkaline solutions have more hydroxide ions (higher than 7 on the pH scale).
• Logarithmic scale: each pH number change represents a tenfold change in hydrogen ion concentration.
  • pH of 7 = 10^{-7} moles of hydrogen ions per liter.
  • pH of 6 is 10 times greater concentration of hydrogen ions than pH of 7.
  • pH of 2 = 0.01 moles per liter, 10,000 times stronger than pH of 6.

Examples

• Solution with a pH of 5: hydrogen ion concentration is 10^{-5}.
• A solution of pH 4 has 1,000 times greater hydrogen ion concentration than a solution of pH 7.
• pH of 11 has a lower hydrogen ion concentration than a pH of 9.
• Small pH adjustments can make significant differences due to the scale's exponential nature.
• Hand soap is extremely alkaline (pH of 10), harsh on skin compared to skin's pH of 5.5 - 5.6.
• Surfactant formulas (shampoos, body washes) are pH-balanced to be closer to skin's pH.
• Exposure to water (pH 7) can cause dermatitis due to the difference in hydrogen ion concentration compared to skin.
• Orange juice is a little bit more like our peel sort of pH, for example.

Strong and Weak Acids

• Inorganic acids, bases, and salts dissolve and dissociate into ions in water.
• Water molecules form a hydration shell around the ions.
• Acids: Substances that donate hydrogen ions to the solution when dissolved, increasing hydrogen ion concentration.
• Referred to as proton donors.
• Hydrogen ions attach to water molecules, forming hydronium ions.
• Strong acids dissociate completely and irreversibly.
  • Hydrochloric acid (HCl) dissociates fully into hydrogen ions and chloride ions.
  • Sulfuric acid has solutions that have high concentrations of high hydrogen ions when it's placed into an aqueous solution or into water.
• Weak acids, such as carbonic acid and acetic acid (ethanoic acid), can reform initial reactants and only partially dissociate.
• Dissociation constant (pKa) measures the level of dissociation. Stronger acids have lower pKa values.

pKa value

• Lower pKa values indicate stronger acids, requiring more caution.
• Higher pKa values indicate gentler acids with fewer complications.
• Logarithmic scale ranging from 0.12 to 52.
• Strong acids values are closer to zero
• Weaker acids have values closer to nine
  • Acids with pKa values less than 3 require extra caution.
• Hydrochloric acid: pKa = 0.7
• Acetic acid (vinegar): pKa = 4.75
• Glycolic acid: pKa = 3.83
• Lactic acid: pKa = 3.86
• Pyruvic acid: pKa = 2.49
• Trichloroacetic acid (TCA): pKa = 0.26 (very strong).

Bases

• Reduce hydrogen ion concentration in a solution.
• Two mechanisms:
  • Directly accepting hydrogen ions.
  • Increasing hydroxide ion concentration.
• Some bases (e.g., ammonia) attract hydrogen ions, forming ammonium ions.
• Other bases (e.g., sodium hydroxide) dissociate in solution, increasing hydroxide ion levels.
• Reactions with two-directional arrows are weaker bases.
• Unidirectional arrows indicate stronger bases like sodium hydroxide, which dissociates almost completely.

Salts and Neutralization

• Mixing acids and bases can form salts and water.
• Example: Hydrogen chloride and sodium hydroxide dissociate, then rearrange to form water and sodium chloride (table salt).
• When hydrogen chloride and sodium hydroxide are added into water, the hydrogen and the chloride ions separate. In the case of sodium hydroxide, sodium and hydroxide ions separate out.
• Then there's rearrangement and reforming of those.

Buffers

• Aqueous solutions containing a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.
• Keep pH at a specific range.
• Internal buffering systems maintain narrow pH limits in the body.
• Blood pH: 7.35 to 7.45 (slightly alkaline).
• Saliva: Slightly alkaline.
• Interstitial fluid: slightly alkaline.
• Intracellular: Near neutral.
• Gastric juices: Low pH of 1 to 3 (high hydrogen ion concentration for digestion).
• Compatibility with life: pH ranges from 6.8 to 7.9.
• Skin's acidic pH is a defense against microbes.
• Respiratory and renal systems, along with circulating buffer systems, adjust pH levels.
• Carbonic acid-bicarbonate buffer system: Carbonic acid (H2CO3) dissociates into hydrogen ions and carboxyl ions; bicarbonate molecule (HCO3) acts as a weak base.

Buffers in Peels

• Keep pH of a solution constant, ensuring stability.
• Buffered peels are safer and easier to use.
• Free acid peels require more caution.
• Neutralization: Balancing an acid by adding an alkaline substance to reach a pH of 7.
• Lactic acid can be neutralized by sodium bicarbonate solution.
• Sodium bicarbonate solution is used to remove those extra hydrogen ions.
• Self-neutralizing peels: Salicylic acid (BHA), trichloroacetic acid, Jessner peels, phenol peels; excess is removed, rather than neutralizing the peels.
• Water is enough to remove the excess of the peels.
• Alpha hydroxy acid peels need to be neutralized at the end of the procedure. Neutralizing Solution
  • Is needed as a safety precaution measures in case of frosting in the skin due to damage to skin tissues.
• Partially neutralized peels: Acid mixed with a small amount of base, leading to a gradual release of free acid.
• Gives more control to this process, and makes the product safer.